1. Atomic Structure, Periodicity, and Bonding

2. Sizes of Ions Ionic size depends upon: The nuclear charge.
The number of electrons.
The orbitals in which electrons reside.
 2009, Prentice-Hall, Inc.

3. Sizes of Ions Cations are smaller than their parent atoms.
The outermost electron is removed and repulsions between electrons are reduced.
 2009, Prentice-Hall, Inc.

4. Atomic Radii & Size of Ions
Atoms get smaller left to right because shielding remains roughly constant as nuclear charge increases across a period.
Atoms get larger top to bottom because outermost electrons have higher principal quantum numbers. They are farther away from the nucleus.
Cations get smaller b/c
e- in outermost shells are gone
Fewer e- to e- repulsions
Anions get bigger b/c
additional e- have increased repulsions causing the need for more space.

5. Sizes of Ions Ions increase in size as you go down a column.
This is due to increasing value of n.
 2009, Prentice-Hall, Inc.

6. Sizes of Ions
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.
 2009, Prentice-Hall, Inc.

7. Quantum Numbers

8. Shielding
Effective nuclear charge (Zeff) is the charge experienced by an electron on a many-electron atom. The effective nuclear charge is not the same as the charge on the nucleus because of the effect of inner electrons. An electron is attracted to the nucleus, but repelled by electrons that shield or screen it from the full nuclear charge. This shielding is called the screening effect, OR THE SHIELDING EFFECT! The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of
electrons that block it from the nucleus

9. Electron Domains
We can refer to the electron pairs as electron domains.
In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain.
The central atom in this molecule, A, has four electron domains.
 2009, Prentice-Hall, Inc.

10. Nonbonding Pairs and Bond Angle
Nonbonding pairs are physically larger than bonding pairs.
Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.
 2009, Prentice-Hall, Inc.

11. Larger Molecules
In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole.
 2009, Prentice-Hall, Inc.

12. Larger Molecules
This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.
 2009, Prentice-Hall, Inc.

13. Overlap and Bonding
We think of covalent bonds forming through the sharing of electrons by adjacent atoms.
In such an approach this can only occur when orbitals on the two atoms overlap.
 2009, Prentice-Hall, Inc.

14. Overlap and Bonding
Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron-electron repulsion.
However, if atoms get too close, the internuclear repulsion greatly raises the energy.
 2009, Prentice-Hall, Inc.

15. Hybrid Orbitals
But it’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize.
 2009, Prentice-Hall, Inc.

16. Hybrid Orbitals Consider beryllium:
In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.
 2009, Prentice-Hall, Inc.

17. Hybrid Orbitals
But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.
 2009, Prentice-Hall, Inc.

18. Hybrid Orbitals
Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.
These sp hybrid orbitals have two lobes like a p orbital.
One of the lobes is larger and more rounded as is the s orbital.
 2009, Prentice-Hall, Inc.

19. Hybrid Orbitals
These two degenerate orbitals would align themselves 180 from each other.
This is consistent with the observed geometry of beryllium compounds: linear.
 2009, Prentice-Hall, Inc.

20. Hybrid Orbitals
With hybrid orbitals the orbital diagram for beryllium would look like this.
The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.
 2009, Prentice-Hall, Inc.

21. Hybrid Orbitals Using a similar model for boron leads to…
 2009, Prentice-Hall, Inc.

22. Hybrid Orbitals …three degenerate sp2 orbitals.
 2009, Prentice-Hall, Inc.

23. Hybrid Orbitals
With carbon we get…
 2009, Prentice-Hall, Inc.

24. Hybrid Orbitals …four degenerate sp3 orbitals.
 2009, Prentice-Hall, Inc.

25. Hybrid Orbitals
For geometries involving expanded octets on the central atom, we must use d orbitals in our hybrids.
 2009, Prentice-Hall, Inc.

26. Hybrid Orbitals This leads to five degenerate sp3d orbitals…
…or six degenerate sp3d2 orbitals.
 2009, Prentice-Hall, Inc.

27. Hybrid Orbitals
Once you know the electron-domain geometry, you know the hybridization state of the atom.
 2009, Prentice-Hall, Inc.

28. Valence Bond Theory
Hybridization is a major player in this approach to bonding.
There are two ways orbitals can overlap to form bonds between atoms.
 2009, Prentice-Hall, Inc.

29. Sigma () Bonds Sigma bonds are characterized by Head-to-head overlap.
Cylindrical symmetry of electron density about the internuclear axis.
 2009, Prentice-Hall, Inc.

30. Pi () Bonds Pi bonds are characterized by Side-to-side overlap.
Electron density above and below the internuclear axis.
 2009, Prentice-Hall, Inc.

31. Single Bonds
Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering.
 2009, Prentice-Hall, Inc.

32. Multiple Bonds
In a multiple bond one of the bonds is a  bond and the rest are  bonds.
 2009, Prentice-Hall, Inc.

33. Multiple Bonds
In a molecule like formaldehyde (shown at left) an sp2 orbital on carbon overlaps in  fashion with the corresponding orbital on the oxygen.
The unhybridized p orbitals overlap in  fashion.
 2009, Prentice-Hall, Inc.

34. Multiple Bonds
In triple bonds, as in acetylene, two sp orbitals form a  bond between the carbons, and two pairs of p orbitals overlap in  fashion to form the two  bonds.
 2009, Prentice-Hall, Inc.

35. Delocalized Electrons: Resonance
When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.
 2009, Prentice-Hall, Inc.

36. Delocalized Electrons: Resonance
In reality, each of the four atoms in the nitrate ion has a p orbital.
The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.
 2009, Prentice-Hall, Inc.

37. Delocalized Electrons: Resonance
This means the  electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.
 2009, Prentice-Hall, Inc.

38. Resonance
The organic molecule benzene has six  bonds and a p orbital on each carbon atom.
 2009, Prentice-Hall, Inc.

39. Resonance
In reality the  electrons in benzene are not localized, but delocalized.
The even distribution of the electrons in benzene makes the molecule unusually stable.
 2009, Prentice-Hall, Inc.